Question

You know that an unlabeled bottle contains an aqueous solution of one of the following: \(\mathrm{AgNO}_{3}, \mathrm{CaCl}_{2}\), or \(\mathrm{Al}_{2}\left(\mathrm{SO}_{4}\right)_{3}\). A friend suggests that you test a portion of the solution with \(\mathrm{Ba}\left(\mathrm{NO}_{3}\right)_{2}\) and then with \(\mathrm{NaCl}\) solutions. According to your friend's logic, which of these chemical reactions could occur, thus helping you identify the solution in the bottle? (a) Barium sulfate could precipitate. (b) Silver chloride could precipitate. (c) Silver sulfate could precipitate. (d) More than one, but not all, of the reactions described in answers a-c could occur. (e) All three reactions described in answers a-c could occur.

Short answer

(d) More than one, but not all, of the reactions described in answers a-c could occur.

Step-by-step solution

Investigate the reactions of each solution with Ba(NO3)2

To investigate the possible reactions, we must look at the potential reactions by combining the compounds in the solutions: 1. Reaction with \(\mathrm{AgNO}_{3}\): The possible product is \(\mathrm{Ba}\left(\mathrm{NO}_{3}\right)_{2}\) + \(\mathrm{AgNO}_{3}\) → \(\mathrm{BaAg}\left(\mathrm{NO}_{3}\right)_{3}\), which is not a precipitate. So, no reaction can take place for option (a) or (b), but (c) might be a possibility because we don't know if \(\mathrm{NaCl}\) is present. 2. Reaction with \(\mathrm{CaCl}_{2}\): The possible product is \(\mathrm{Ba}\left(\mathrm{NO}_{3}\right)_{2}\) + \(\mathrm{CaCl}_{2}\) → \(\mathrm{BaCl}_{2}\) + \(\mathrm{Ca}\left(\mathrm{NO}_{3}\right)_{2}\). Barium chloride, \(\mathrm{BaCl}_{2}\), is soluble, so no reaction can form a precipitate. 3. Reaction with \(\mathrm{Al}_{2}\left(\mathrm{SO}_{4}\right)_{3}\): The possible product is \(\mathrm{Ba}\left(\mathrm{NO}_{3}\right)_{2}\) + \(\mathrm{Al}_{2}\left(\mathrm{SO}_{4}\right)_{3}\) → \(\mathrm{BaSO}_{4}\) + \(\mathrm{Al}\left(\mathrm{NO}_{3}\right)_{3}\). Barium sulfate, \(\mathrm{BaSO}_{4}\), is insoluble and forms a precipitate, which proves the possibility of option (a).

Investigate the reactions of each solution with NaCl

Now we will look at the potential reactions when we add \(\mathrm{NaCl}\) to each of the solutions: 1. Reaction with \(\mathrm{AgNO}_{3}\): The possible product is \(\mathrm{AgNO}_{3}\) + \(\mathrm{NaCl}\) → \(\mathrm{AgCl}\) + \(\mathrm{NaNO}_{3}\). Silver chloride, \(\mathrm{AgCl}\), is insoluble and forms a precipitate, which proves the possibility of option (b). 2. Reaction with \(\mathrm{CaCl}_{2}\): The possible product is \(\mathrm{CaCl}_{2}\) + \(\mathrm{NaCl}\), which doesn't show any potential precipitation of insoluble compounds. 3. Reaction with \(\mathrm{Al}_{2}\left(\mathrm{SO}_{4}\right)_{3}\): The possible product is \(\mathrm{Al}_{2}\left(\mathrm{SO}_{4}\right)_{3}\) + \(\mathrm{NaCl}\), which doesn't show any potential precipitation of insoluble compounds.

Conclusion

Based on the investigation of the reactions with both \(\mathrm{Ba}\left(\mathrm{NO}_{3}\right)_{2}\) and \(\mathrm{NaCl}\), we can conclude the following chemical reactions could occur: (a) Barium sulfate could precipitate (reaction with \(\mathrm{Al}_{2}\left(\mathrm{SO}_{4}\right)_{3}\)). (b) Silver chloride could precipitate (reaction with \(\mathrm{AgNO}_{3}\)). Option (c) was not proven to be possible through any of the combinations, so the correct answer is: (d) More than one, but not all, of the reactions described in answers a-c could occur.

Key concepts

Aqueous Solutions

In chemistry, an aqueous solution is a solution where water serves as the solvent. The term "aqueous" refers to water, derived from the Latin word "aqua". In aqueous solutions, various substances get dissolved in water, creating a homogeneous mixture.

Water is often called the "universal solvent" because it can dissolve a wide range of compounds. The reason for this is its polar nature, which allows it to interact with various ions and molecules, pulling them apart and evenly distributing them throughout the solution.

Aqueous solutions are crucial in many chemical reactions because they allow reactants to mix thoroughly. This thorough mixing is necessary for reactions to occur quickly and completely. Additionally, when ionic compounds dissolve in water, they disassociate completely into their ions, which can readily react with other ions. For example, when \ \( \mathrm{AgNO}_{3} \ \) dissolves in water, it dissociates into \ \( \mathrm{Ag}^{+} \ \) and \ \( \mathrm{NO}_{3}^{-} \ \) ions. This disassociation is often key to precipitation reactions, where newly formed ions may combine to form insoluble compounds that precipitate out of solution.

• Aqueous solutions are central to many solubility tests.
• They facilitate effective mixing of reagents.
• Water's polar nature is ideal for dissolving ionic substances.

Precipitation Reactions

Precipitation reactions occur when two aqueous solutions combine to form an insoluble product known as a precipitate. These reactions are a type of double displacement reaction, where ions in the reactant solutions switch partners to form new compounds, one of which is insoluble in water.

The formation of a solid precipitate is often visible as a cloudiness or with particles settling out of the solution. This solid can be separated from the liquid solution by filtration. For a precipitation reaction to take place, the product formed must be insoluble in the aqueous medium.

An example from the original exercise involves mixing an aqueous solution of \ \( \mathrm{AgNO}_{3} \ \) with \ \( \mathrm{NaCl} \ \), resulting in the formation of silver chloride, \ \( \mathrm{AgCl} \ \), a common precipitate in solubility tests. Similarly, \ \( \mathrm{BaSO}_{4} \ \) can precipitate when \ \( \mathrm{Ba(NO}_{3})_{2} \ \) is added to \ \( \mathrm{Al}_{2}(\mathrm{SO}_{4})_{3} \ \).

Key aspects of precipitation reactions include:

• The need for at least one insoluble product to form.
• Visual identification of precipitates makes them useful in analytical chemistry.
• The reaction's ability to test for the presence of certain ions.

Solubility Rules

Solubility rules are guidelines that help predict whether a particular compound will dissolve in water or form a precipitate. These rules are based on empirical observations of many compounds' behavior in aqueous solutions. Solubility is a key factor in determining whether a reaction will result in precipitation.

Here are some basic solubility rules:

• Most nitrate salts (e.g., \ \( \mathrm{NO}_{3}^{-} \ \)) are soluble in water.
• Salts of alkali metals (such as sodium, \ \( \mathrm{Na}^{+} \ \)) and ammonium (\